The  Llectrical  Conductivity 

of 
Certain  Salts  In  Pyridine 


THL5I5 


SUBMITTED  TO  THE  FACULTY  OF  THE  GRADUATE  COLUCGK 

OF  THE  STATE  UNIVERSITY  OF  IOWA  IN  PARTIAL 

FULFILLMENT  OF  THE   REQUIREMENTS 

FOR  THE  DEGREE  OF  DOCTOR 

OF  PHILOSOPHY 


BY 

EDWARD  X.  ANDERSON 

IOWA  CITY,  IOWA 


EASTON,  PA.: 

ESCHENBACH  PRINTING  Co 
1915 


The  Electrical  Conductivity 

of 
Certain  Salts  In  Pyridine 


THL5I5 


SUBMITTED  TO  THE  FACULTY  OF  THE  GRADUATE  COLLEGE 

OF  THE  STATE  UNIVERSITY  OF  IOWA  IN  PARTIAL 

FULFILLMENT  OF  THE  REQUIREMENTS 

FOR  THE  DEGREE  OF  DOCTOR 

OF  PHILOSOPHY 


BY 

EDWARD  X.  ANDERSON 

IOWA  CITY,  IOWA 


EASTON,  PA.: 

ESCHENBACH  PRINTING  Co 
1915 


So 


«* 


THE    ELECTRICAL    CONDUCTIVITY    OF    CERTAIN 
SALTS  IN  PYRIDINE 


BY   EDWARD   X.    ANDERSON 

Introduction 

Thus  far  little  work  of  a  systematic  nature  has  been 
done  on  the  electrical  conductivity  of  salts  in  pyridine,  and 
this  has  been  limited  in  every  case  to  one  temperature  and  to 
the  more  dilute  solutions.  Laszczynski  and  Gorski1  inves- 
tigated the  conductivity  of  pyridine  solutions  containing 
lithium  chloride,  and  the  iodides  and  thiocyanates  of  potas- 
sium, sodium  and  ammonium,  all  measurements  being  made 
at  1 8°.  They  found  that  the  equivalent  curves  showed 
sufficient  convergence  to  render  possible  the  calculation  of 
/QQ  by  extrapolation,  where  /^  represents  the  equivalent  elec- 
trical conductivity  of  the  solutions  at  infinite  dilution.  The 
maximum  values  of  the  equivalent  conductivity  lie  between 
40  and  46,  and  they  are,  therefore,  as  large  as  the  values  ob- 
tained, using  ethyl  alcohol  as  a  solvent.  Lithium  chloride 
showed  scarcely  any  dissociation.  Lincoln2  found  that  various 
inorganic  salts  yielded  conducting  solutions  in  pyridine,  but 
values  for  /^  are  rarely  encountered.  Dutoit  and  Duper- 
thuis3  determined  the  conductivities  of  pyridine  solutions 
containing  potassium  iodide,  potassium  thiocyanate  and 
sodium  thiocyanate,  for  dilutions  ranging  from  1,000  to 
20,000  liters  per  gram-mole  of  salt.  Within  these  limits 
Ostwald's  dilution  law  applies  in  most  cases,  thus  rendering 
it  possible  to  calculate  the  degree  of  dissociation  at  any  given 
dilution. 

Jones  and  West4  studied  the  temperature  coefficients 
of  conductivity  in  aqueous  solutions  and  the  effect  of  tem- 
perature on  dissociation.  In  their  work  the  following  rela- 

1  Zeit.  Elektrochemie,  4,  290-293  (1897). 

2  Jour.  Phys.  Chem.,  3,  457-484  (1899): 

3  Jour.  chim.  phys.,  6,  699-725  (1909). 

4  Am.  Chem.  Jour.,  34,  357  (1905);  35.  445  (1906). 


754  Edward  X.  Anderson 

lions  were  found  to  exist:  a  rise  in  temperature,  ranging  from 
o°  to  35°  produces  a  large  increase  in  conductivity  due  to  an 
increase  in  ionic  mobility.  This  last  effect  results  from  a 
reduction  in  viscosity  and  a  simplification  of  complexes. 
For  any  given  electrolyte  the  temperature  coefficients  of  con- 
ductivity increase  with  increasing  dilution;  with  different 
electrolytes  this  increase  is  greatest  for  those  electrolytes 
with  large  hydrating  power.  Voellmer1  found  that  the 
temperature  coefficients  of  electrical  conductivity  increased 
with  rising  dilution  for  solutions  of  lithium  chloride,  and  the 
acetates  and  iodides  of  potassium  and  sodium  in  methyl 
alcohol,  and  also  for  the  solutions  of  sodium  chloride,  calcium 
chloride,  silver  nitrate  and  calcium  nitrate  in  ethyl  alcohol. 
The  molecular  conductivities  of  various  concentrations  of 
the  chlorides  of  nickel,  manganese  and  cobalt  in  methyl  alco- 
hol, ethyl  alcohol  and  acetone,  between  o°  and  45°,  were 
studied  by  Rimbach  and  Weitzel.2  In  contrast  with  the 
near  constancy  of  the  temperature  coefficients  on  dilution  in 
aqueous  solutions,  they  found  that  in  organic  solvents  the 
temperature  coefficients  increase,  as  a  rule,  with  increasing 
dilution.  An  examination  of  the  work  of  Jones  and  Clover,3 
and  of  Jones  and  West,4  however,  shows  that  the  tempera- 
ture coefficients  of  conductivity  in  aqueous  solutions  do  in- 
crease with  increase  in  dilution  and  in  most  cases  the  increase 
is  quite  pronounced. 

The  results  obtained  and  the  conclusions  made  by  the 
above  investigators  will  be  considered  more  fully  in  the  dis- 
cussion of  the  present  work. 

Pyridine,  like  water,  has  the  power  of  combining  with 
salts  to  form  crystalline  solids  with  pyridine  of  crystalliza- 
tion. This  being  the  case,  according  to  the  law  of  mass  ac- 
tion, one  should  expect  these  salts  to  combine  with  a  much 
larger  amount  of  pyridine  when  in  solution  in  this  solvent. 

1  Wied.  Ann.,  52,  328  (1894). 

2  Zeit.  phys.  Chem.,  79,  279  (1912). 

3  Am.  Chem.  Jour.,  43,  187  (1910). 

4  Loc.  cit. 


Electrical  Conductivity  of  Certain  Salts  in  Pyridine      755 

Further,  as  in  the  case  of  hydrates,  one  should  expect  the 
complexity  of  these  ionic  ' '  pyridinates "  to  be  greatest  in  the 
most  dilute  solutions,  and  also  that  the  complexity  of  the 
solvate  'at  any  given  dilution  should  decrease  with  rise  in 
temperature.  This  is  based  on  the  general  law  that,  the 
stability  of  complexes  decreases  with  the  rise  in  temperature. 
Since  pyridine  tends  to  form  these  "pyridinates,"  interesting 
results  might  well  be  expected  which,  in  a  certain  sense,  are 
parallel  to  those  obtained  in  aqueous  solutions. 

It  was,  therefore,  thought  worth  while  to  make  a  care- 
ful, systematic  study  of  the  conductivity  of  solutions  of 
various  salts  in  pyridine  at  different  temperatures  and  over 
as  wide  a  range  of  concentrations  as  the  experimental  condi- 
tions would  permit. 

Experimental 

For  measuring  the  conductivity  the  well  known  Kohl- 
rausch  method,  consisting  of  the  Wheatstone  bridge,  induc- 
tion coil  and  telephone  receiver,  was  used.  The  resistance 
boxes  were  certified  by  the  Reichsanstalt  and  the  Bureau  of 
Standards. 

The  conductivity  cells,  three  in  number,  were  of  the 
type  first  used  by  Jones  and  Lindsay.1  These  were  pro- 
vided with  ground  glass  stoppers  and  sealed-in  electrodes. 
A  o .  02  N  potassium  chloride  solution  was  used,  in  determining 
the  cell  constants.  The  "chemically  pure"  potassium  chlor- 
ide was  recrystallized  from  conductivity  water  and  carefully 
dried  by  heating  for  some  time  at  a  dull  red  heat  and  then 
cooled  in  a  desiccator.  The  salt  was  then  dissolved  in  con- 
ductivity water  prepared  by  the  method  of  Jones  and  Mackay.* 
This  water  had  a  specific  conductance  of  i  .5-  io~6  r.  o.  at  25°. 
The  specific  conductances  of  the  0.02  N  potassium  chloride 
at  o°  and  25°  were  taken  as  0.001522  and  0.002768,  respec- 
tively.3 The  cell  constants  thus  determined  were  then 
checked  against  0.002  N  potassium  chloride.  One  set  of 

1  Zeit.  phys.  Chem.,  56,  129  (1906). 

2  Am.  Chem.  Jour.,  19,  90  (1897). 

3  Ostwald-Luther:  Messungen,  third  ed.,  474 


756  Edward  X.  Anderson 

constants  used  for  the  three  cells  were  0.0438,  0.2176  and 
0.3985,  respectively,  and  the  values  of  lv  for  0.002  N  potas- 
sium chloride  checked  to  within  0.3  percent. 

" Chemically  pure"  pyridine  was  allowed  to  stand  over 
fused  caustic  potash  for  several  months.  It  was  then  de- 
canted and  distilled.  The  fraction  passing  over  between 
115°  and  116.1°  at  74.5  cms  was  collected,  the  first  and  last 
portions  being  rejected.  The  specific  conductance  at  o° 
was  found  to  be  o.o57-io~7  r.  o.,  at  25°,  o.74-io~7  and  at  50°, 
i.2-io~7.  Lincoln1  found  the  specific  conductance  of  this 
pyridine  to  be  7.6-  io~7  at  25°,  a  much  larger  value. 

Unless  otherwise  indicated,  it  will  be  understood  that 
Kahlbaum's  best  grade  "C.  P."  chemicals  were  used.  After 
complete  dehydration,  by  methods  to  be  mentioned  later, 
the  salts  were  preserved  in  tightly  stoppered  weighing  bottles 
over  phosphorus  pentoxide.  Trial  solubility  tests  were  made 
to  determine  approximately  the  amount  of  each  salt  required 
for  saturation  at  room  temperature.  Those  salts  which  do 
not  show  hygroscopic  properties  were  weighed  out  directly 
and  the  exact  amounts  required  for  a  normality  conveniently 
close  to  that  of  complete  saturation  was  taken,  but  for  those 
salts  which  tend  to  absorb  moisture  the  method  of  weighing 
by  difference  was  used.  In  diluting  the  mother  solutions 
every  possible  care  was  taken  to  prevent  contact  with  mois- 
ture. Suction  was  applied  directly  to  the  pipette  through  a 
calcium  chloride  drying  tube. 

All  the  solutions  were  prepared  at  room  temperatures, 
2i°-22°,  and  transferred  to  tightly  fitting  glass-stoppered 
bottles.  The  conductivity  measurements  were  made  as  soon 
after  preparation  as  possible. 

The  capacities  of  the  25°  and  50°  baths  used  were  20 
and  1 6  liters,  respectively.  They  were  heated  by  means  of 
immersed  electric  lights,  which  were  automatically  operated 
by  means  of  a  contact  toluol  regulator  and  relay  system. 
Stirrers,  propelled  by  a  small  motor,  were  used  to  agitate  the 

1  Jour.  Phys.  Chem.,  3,  457  (1899). 


Electrical  Conductivity  of  Certain  Salts  in  Pyridine      757 

water  vigorously  and  thereby  maintain  a  uniform  tempera- 
ture throughout  the  baths.  By  this  means  both  tempera- 
tures were  automatically  kept  at  25°  ±  0.02  and  50°  =•=  0.05. 
A  mixture  of  well  washed,  finely  crushed  ice  moistened  with 
distilled  water  was  used  for  the  o°  bath.  The  order  of  tem- 
peratures followed  for  the  conductivity  measurements  was 
25°,  o°  and  50°.  The  temperature  25°  was  selected  as  the 
initial,  because  the  solutions  were  made  up  at  room  tempera- 
ture. Thermal  equilibrium  would,  therefore,  be  more  quickly 
attained  at  this  temperature  than  at  either  of  the  others. 
The  temperature  o°  was  selected  as  the  second  in  order  to 
permit  of  the  minimum  change  in  the  concentration  of  the 
solutions  by  condensation  of  pyridine  on  the  inside  wall  of 
the  cells.  If  the  cells  had  been  immersed  in  the  50°  bath 
before  being  subjected  to  a  temperature  of  o°,  there  would 
be  more  pyridine  condensed  on  the  cell  walls  than  if  the 
25°  immediately  preceded  the  o°  bath.  Forty  minutes  were 
found  to  be  sufficient  for  establishing  thermal  equilibrium, 
because  by  trial  experiments  the  conductivity  was  found  to 
remain  constant  after  this  amount  of  time  had  elapsed. 

Silver  Nitrate. — The  pure  crystals  were  pulverized  and 
kept  in  the  dark  over  phosphorus  pent  oxide  for  several  days. 
This  salt  dissolves  in  pyridine  with  a  considerable  evolution 
of  heat,  probably  due  to  the  formation  of  AgNO3-2Pyr,  or 
AgNO3«3Pyr,  both  of  which  have  been  found  capable  of 
existing  in  the  solid  phase.1  Kahlenberg  and  Brewer2  found 
AgNO3-3Pyr  to  be  stable  between  — 24°  and  +48.5°,  where 
it  changes  to  AgNO3-2Pyr.  Below  —24°  the  compound 
AgNO3-6Pyr  is  the  stable  form. 

In  the  following  data  V  denotes  the  volume  of  the  solu- 
tion in  liters  containing  one  gram-equivalent  weight  of  the 
salt.  IQ,  /25  and  Ib0  represent  the  equivalent  conductivities 
of  the  various  solutions  at  o°,  25°  and  50°,  respectively. 
The  temperature  coefficients  are  headed  by  the  letters  A,  B 
and  C  where 


1  Reitzenstein:  Liebig's  Ann.,  282,  267  (1894). 

2  Jour.  Phys.  Chem.,  12,  283  (1908). 


758 


Edward  X.  Anderson 


A  = 


/25 


B 


j  ^ 
and  C  = 


/o  •  25  '  /25  •  25  /0  •  50  ' 

In  the  curves  which  follow,  the  cube  roots  of  the  volumes 
in  cubic  centimeters  are  plotted  as  abscissae,  against  the 
equivalent  conductivity  I  as  ordinates.  Each  space  along 
the  abscissa  is  equal  to  2 . 5  units  after  the  root  is  extracted. 
Every  space  along  the  ordinate  is  equal  to  10  units  after  the 
equivalent  conductivity  is  multiplied  by  the  factor,  which 
will  be  specified  in  each  case.  This  factor  is  used  in  order 
to  magnify  the  trend  of  the  curves.  It  must  also  be  remem- 
bered that  any  errors  are  likewise  highly  magnified. 

TABLE  I 


V 

k 

/25 

/50 

I 

1.05 

i-55 

2.01 

2 

14-77 

19.38 

23.28 

IO 

20.68 

25.38 

27.25 

20 

22.38 

27.05 

29.17 

ioo         27.80 

34-49 

37-92 

500             37-31 

47-63 

55-iQ 

TABLE  II 


V 

A 

B 

C 

I 

0.0149 

O.OI2I 

0.0183 

2 

0.0125 

O.OOSl 

0.0115 

10 

0.0091 

O.OO29 

o  .  0064 

2O 

o  .  0084 

o  .  003  i 

o  .  006  i 

IOO 

o  .  0096 

o  .  0040 

o  .  0073 

500 

O.OIII 

0.0063 

o  .  0095 

From  Fig.  I  it  will  be  seen  that,  with  dilution  the  equiva- 
lent conductivity  of  silver  nitrate  increases  at  first  very  rapidly 
and  then  less  rapidly  with  further  dilution  for  alt  three  tem- 
peratures. They  do  not  appear  to  approach  maximum  values. 
The  values  for  lv  here  given  agree  very  closely  with  those 
given  by  Lincoln1  for  the  same  salt  at  25°.  The  tempera- 


1  Loc.  cit. 


Electrical  Conductivity  of  Certain  Salts  in  Pyridine      759 

ture  coefficients  (Table  II)  show  distinct  minima,  decreasing 
at  first  very  rapidly  and  then  increasing  very  slowly  on  further 
dilution.  Although  solutions  of  silver  nitrate  in  pyridine 
possess  a  relatively  high  molecular  conductivity,  Walden 
and  Centnerszwer1  found  that  the  molecular  weights  of  sil- 
ver nitrate  in  dilute  pyridine  solutions  are  normal,  while 


Fig.  I 

in  concentrated  solutions  (o .  i  N  to  i .  o  N)  the  molecular 
weights  are  greater  than  normal,  thus  indicating  association. 
By  the  same  method  Schmuilow2  found  that  this  salt  is  ap- 
parently non -ionized.  Since  transference  measurements  made 
by  Neustadt  and  Abegg3  showed  that  both  Ag+  ions  and  NO3 
radicle  migrated  toward  the  cathode,  they  assumed  that,  if 
ionization  does  take  place,  it  does  so  according  to  the  ex- 


pression 
d) 


(AgN03)2  ±1^  Ag2+N03  +  NO,-. 


1  Zeit.  phys.  Chem,  55,  321  (1906). 

2  Zeit.  anorg.  Chem.,  15,  18  (1897). 

3  Zeit.  phys.  Chem.,  69,  486  (1910). 


760 


Edward  X.  Anderson 


That  simple^Ag+  ions  are  also  present  to  a  slight  extent  is 
not  to  be  doubted,  and,  therefore,  the  more  complete  equi- 
librium may  be  represented  by  the  following  equation: 
(2)  2Ag+  +  2NO3-±^:2AgNO3±^:(AgNO3)2^^:Ag2+NO3  +  NO3~. 

Lithium  Chloride. — The  sample  was  heated  at  120°  for 
several  days,  with  frequent  pulverizing  in  a  hot  agate  mor- 
tar until  the  tendency  to  cake  ceased.  The  dry,  finely  pow- 
dered salt  was  preserved  in  a  tightly  stoppered  weighing 
bottle. 

Lithium  chloride  dissolves  in  pyridine  with  a  evolu- 
tion of  heat;  it  separates  from  solution  as 

TABLE  III 


V 

k 

*» 

/60 

0.59 

0.143 

0.199 

0.239 

I  .00 

0.218 

0.264 

0.282 

2  .OO 

0.254 

0.290 

0.299 

10.00 

0.279 

0.322 

0.346 

100.00 

0-519 

0.573 

0.613 

IOOO.OO 

1.47 

1  .60 

1.68 

TABLE  IV 


V 

A 

B 

c 

0-59 

0.0160 

o  .  0079 

0.0135 

I  .00 

0.0083 

0.0028 

0.0058 

2.00 

o  .  0056 

0.0012 

o  .  0035 

IO.OO 

0.0061 

o  .  0030 

o  .  0047 

100.00 

0.0041 

0.0028 

o  .  0036 

1000.00 

0.0037 

O.OO2O          O.OO29 

An  examination  of  Table  III  and  Fig.  II  shows  that 
lithium  chloride  is  at  best  a  very  poor  conductor  and  is  but 
slightly  dissociated  at  all  concentrations  and  temperatures. 
The  equivalent  conductivities  do,  however,  show  a  rapid 
increase  at  first,  and  then  very  slight  and  finally  more  rapidly, 
with  increasing  dilution.  The  values  found  by  Laszczynski 
and  Gorski2  for  the  same  solutions  are  about  four  times  as- 

1  Laszczynski:  Ber.  deutsch.  chem.  Ges.,  27,  2285  (1894). 

2  Loc.  cit. 


Electrical  Conductivity  of  Certain  Salts  in  Pyridine      761 

large,  due,  perhaps,  to  the  presence  of  a  slight  trace  of  mois- 
ture.    Increasing  dilution  has  a  marked  intitial  effect  on  the 


Fig.  II 

temperature  coefficients  (Table  IV),  the  latter  passing  through 
a  minimum. 

Lithium    Bromide. — The    anhydrous    salt    was    prepared 
in  a  manner  similar  to  that  used  for  lithium  chloride. 

TABLE  V 


V 

lo 

fei 

/SO 

0.98 

— 

1.29 

1.65 

2.00 

0.981 

1.72 

1.98 

IO.OO 

2.29 

2.44 

2.40 

100.00 

5-43 

5-34 

4.89 

IOOO.OO 

13.68 

14.  15 

13.58 

10,000.00 

24.8 

28.7 

29.9 

00               (28.5) 

(36.3) 

(49.0) 

TABLE  VI 

V 

A 

B 

c 

0.98 

— 

0.0109 



2.00 

O.O2981                0.006  i 

0.0202 

10.00 

0.0026 

—  0.0006 

0.0009 

100.00 

—  o  .  0007       j       —  o  .  0034 

0.0020 

1000.00 

0.0013 

—  0.0016 

—  o.oooi 

10,000.00 

0.0063 

0.0017 

O.OO4I 

A  glance  at  Fig.  Ill  reveals  the  fact  that  the  equivalent 
conductivity  of  lithium  bromide  increases  steadily  with  in- 
creasing dilutions,  but  not  at  every  dilution  with  a  rise  in 
temperature,  there  being  in  some  cases  a  decrease  in  conduc- 


1  Solidified. 


762 


Edward  X.  Anderson 


tivity  with  a  rise  in  temperature.  The  values  for  lx  at  o°, 
25°  and  50°  were  extrapolated  and  found  to  be  28.5,  36.3 
and  49.0,  respectively.  The  near  proximity  of  the  curves 
signifies  that  a  change  in  temperature  has  little  effect  on  the 
conductivity,  there  being  a  slight  increase  in  this  effect  in 


the  more  dilute  solutions.  The  values  for  the  temperature 
coefficients  (Table  VI)  decrease,  at  first,  with  increasing 
dilution.  This  decrease  is  quite  rapid  in  the  more  concen- 
trated solutions,  passing  through  minima  of  negative  values 
and  then  finally  increasing  at  an  almost  constant  rate. 

Lithium  Iodide. — The  salt  was    carefully  dehydrated  by 
constant  heating  for  several  days  and  nights. 

TABLE  VII 


V 

/o 

ttt 

4o 

1.0         , 

4.40 

7.04 

9.82 

2  .O 

7-79 

10.98 

13.82 

10.  0 

12.76 

16.40 

18.62 

100.  0 

18.34 

23-35 

25.98 

IOOO.O 

27  .  10 

35-99 

42.65 

10,000.0 

31.2 

44-4 

5°  -5 

00 

(31.2) 

(44-9) 

(50-5) 

TABLE  VIII 

V 

A 

B 

C 

I  .0 

0.0224 

0.0158 

0.0246 

2.0 

0.0164 

0.0103 

0.0155 

IO.O 

o.oi  14 

0.0054 

o  .  0092 

IOO.O 

0.0109 

0.0045 

o  .  0083 

IOOO.O 

0.0131 

0.0074? 

0.0115 

10,000.0 

0.0169 

0.0055 

0.0124 

oo 

(0.0180) 

(o  .  0050) 

(0.0124) 

Electrical  Conductivity  of  Certain  Salts  in  Pyridine      763 

From  Fig.  IV  it  is  observed  that  lithium  iodide  is  a  good 
conductor.  The  conductivity  increases  quite  rapidly  at 
the  outset,  but  continuously  increases  at  a  decreasing  rate 
and  soon  attains  maximum  values.  The  o°  and  50°  curves 
appear  to  be  asymptotic  at  a  dilution  of  ten  thousand  liters; 
the  values  for  lv  which  were  actually  obtained  at  this  dilu- 
tion are,  therefore,  taken  as  those  for  lx  and  are  31.2  and 


8 


50.5,  respectively.  By  extrapolation  the  value  for  /^  at 
25°  is  found  to  be  44.9.  The  temperature  coefficients  fall 
exceedingly  rapidly  to  minima  and  then  increase  very  slowly 
with  increasing  dilution. 

Sodium  Iodide. — The  finely  powdered  salt  was  thoroughly 
dried  at  a  temperature  slightly  exceeding  100°.  This  sub- 
stance dissolves  in  pyridine  with  considerable  evolution 
of  heat. 

TABLE  IX 


V 

*• 

b 

/60 

i-33 

0.  II1 

0.70 

0.84 

5.00 

10.00 

ii  .  14 

II  .20 

10.00 

14.56      16.15 

15.80 

100.00 

21.66        23.81 

22.87 

1000.00 

32.99       39-53 

41  .28 

10,000.00 

42.20       56.70 

63.20 

Solid  present. 


764 


Edward  X.  Anderson 
TABLE  X 


V 

A 

B 

c 

i-33 

0.20841 

o  .  0076 

o.  I2791 

5.00 

0.0046 

O  .  OOO2 

o  .  0024 

IO.OO 

o  .  0044 

—  o  .  0009 

0.0017 

100.00 

o  .  0040 

—  -o  .  oo  1  6 

O.OOII 

IOOO.OO 

o  .  0079 

0.0018 

o  .  0050 

10,000.00 

0.0137         0.0046 

o  .  0099 

An  examination  of  Table  IX  and  Fig.  V  reveals  the  fact 
that  lv  follows  the  same  general  trend  as  the  preceding  curves ; 
again,  the  values  of  lv  increase  on  dilution  most  rapidly  in 
the  concentrated  solutions.  By  extrapolation  /^  was  found 
to  be  43.3  at  o°.  Laszczynski  and  Gorski2  obtained  44.32 


5          4 

Fig.  V 


8 


for  the  value  of  /^  at  18°.  For  25°  and  50°,  however,  max- 
imum values  could  not  be  obtained  by  extrapolation.  At 
these  temperatures  the  equivalent  conductivities  continue 
to  increase  with  dilution  more  rapidly  than  at  o°. 

The  temperature  coefficients  show  well  marked  minima. 


1  Solid  phase  present. 

2  Loc.  cit. 


Electrical  Conductivity  of  Certain  Salts  in  Pyridine      765 

Here,    again,    negative   values    appear   for   temperatures   be- 
tween 25°  and  50°. 

Potassium  Thiocyanate. — The  sample  used  was  crys- 
tallized from  absolute  alcohol,  washed  with  absolute  alcohol, 
and  dried  at  95°.  This  salt  differs  from  the  others  studied 
in  that  its  solubility  in  pyridine  decreases  as  the  tempera- 
ture rises. 

TABLE  XI 


V 

/o 

In 

/60 

7.0 

5-97 

7.12 

7-75 

14.0 
70.0 
140.0 
1400.0 
14,000.0 

7.20 
ii  .40 
14.17 
27.32 
42.86 

8.45 

13-36 
16.77 
33-70 

58.51 

9.00 

H-54 
18.14 

38.31 
71.30 

TABLE  XII 

V 

A 

B 

C 

7.0 

14.0 

0.0077 
o  .  0069 

0.0035 

0.0026 

0.0060 
0.0050 

70.0 
140.0 
1400.0 
14,000.0 

o  .  0070       o  .  0035 
o  .  0073       o  .  0033 

0.0093      0.0055 

0.0146        0.0087 

0.0055 
0.0056 
0.0081 
0.0133 

Referring  to  Fig.  VI,  it  is  obvious  that  the  equivalent 
conductivity  of  potassium  thiocyanate  increases  at  a  more 
constant  rate  than  most  of  the  conductivities  of  the  pre- 
ceding salts.  By  extrapolation  46.5  is  obtained  for  the  value 
of  /QQ  at  o°.  Laszczynski  and  Gorski1  give  values  in  fair 
agreement  with  the  above,  tabulated,  comparisons  being  made 
at  1 8°.  They  did  not  obtain  a  value  for  /^  inasmuch  as  the 
dilutions  at  which  they  worked  did  not  exceed  2870.4  liters. 
However,  the  data  which  they  give  show  that  the  conduc- 
tivity is  approaching  a  maximum  value.  The  temperature 
coefficients  exhibit  slight  minima.  In  aqueous  solutions 

1  Loc.  cit. 


Edward  X.  Anderson 


the  effect   of   dilution  upon  the  temperature   coefficients  is 
greatest  in  the  most  dilute  solutions,  as  we  should  expect. 


12545676 
Fig.  VI 

Ammonium  Thiocyanate. — The  anhydrous  salt  was  pre- 
pared in  the  same  manner  as  the  potassium  thiocyanate. 

TABLE  XIII 


V 

/o 

*» 

/SO 

0-33 

2  .  IO 

4.46 

7-43 

i  .00 

8.21 

ii  .70 

15.12 

2  .OO 

10.45 

13.76 

16.53 

10.00 

II  .96 

14.56 

16.29 

100.00 

17  .00 

20.33 

22.  l8 

IOOO.OO 

33-57 

41  .80 

47.76 

TABLE  XIV 


B 


0-33 

1  .00 

2  .OO 
IO.OO 

100.00 
1000.00 


o . 045 i 
0.0170 
0.0127 
0.0087 
0.0078 
o . 0098 


0.0266 
0.0117 
o . 008 i 
o . 0048 
o . 0036 

0.0057 


0.0508 
0.0169 
0.0116 
0.0072 
o . 006 i 
o . 0085 


Electrical  Conductivity  of  Certain  Salts  in  Pyridine      767 

From  Fig.  VII  it  will  be  observed  that  the  equivalent 
conductivity  of  ammonium  thiocyanate  behaves  somewhat 
peculiarly,  when  compared  with  the  previous  curves.  The 
conductivity  curves  rise  rapidly  at  first  and  then  approximate 
parallelism  with  the  volume  axis  in  the  more  concentrated 
solutions,  and  then  increase  again  with  increasing  dilution. 
The  values  for  lv  given  in  the  above  table  are  somewhat 
larger  than  those  given  by  Laszczynski  and  Gorski.  They 
carried  the  dilution  out  to  2080  liters  and  were  able  to  cal- 
culate the  values  for  l^,  which  they  give  as  40.22.  In  the 
above  curves  it  may  be  clearly  seen  that  the  o°  curve  gives 


promise  of  a  maximum  value  for  /„,  but  the  25°  and  50° 
curves  show  practically  no  signs  of  such  a  tendency.  In 
connection  with  the  strange  conduct  of  the  curves  in  the 
vicinity  of  concentrated  solutions,  it  may  be  pointed  out 
that  the  temperature  coefficients  show  equally  striking  changes, 
dropping  suddenly  to  minima,  and  then  increasing  very  slowly 
in  the  dilute  solutions. 

A  relation,  exactly  analogous  to  that  of  the  conduc- 
tivity values  shown  above,  was  observed  by  Franklin1  for 
ammonium  thiocyanate  and  tetramethylammonium  iodide 
in  liquid  sulphur  dioxide.  The  conductivity  curves  of  tetra- 


Jour.  Phys.  Chem.,  15,  675-97  (1911). 


768  Edward  X.  Anderson 

methylammonium  iodide  in  liquid  sulphur  dioxide  are  almost 
indentical  in  form  with  those  here  represented  for  ammonium 
thiocyanate  in  pyridine,  while  those  for  ammonium  thiocyanate 
in  the  same  solvent  are  not  very  unlike  those  obtained  in 
pyridine.  Upon  passing  from  the  most  concentrated  solu- 
tions to  the  most  dilute,  Franklin  found  that  the  molecular 
conductivity  first  increases  to  a  maximum,  then  falls  to  a 
minimum  value,  and  finally  approaches  the  usual  maximum 
on  further  dilution.  He  explained  these  results  by  assuming 
that  the  dissociated  salt  is  auto-ionized  in  the  concentrated 
solutions  and  that  this  effect  decreases  with  dilutions.  On 
the  other  hand,  the  decrease  in  viscosity  with  dilution  causes 
a  rise  in  ionic  mobility.  These  two  effects  balance  each 
other  at  the  first  mentioned  maximum.  On  further  dilution 
the  auto-ionization  disappears  and  the  conductivity  from 
then  on  is  due  to  the  dissociating  power  of  the  solvent.  Con- 
centrated solutions  of  ammonium  thiocyanate  in  pyridine 
are  very  viscous;  the  same  is  true  of  a  concentrated  solution 
of  ammonium  thiocyanate  in  liquid  sulphur  dioxide.1  The 
initial  rapid  increase  in  the  equivalent  conductivities  and  de- 
crease in  temperature  coefficients  are  undoubtedly  chiefly 
due  to  a  rapid  decrease  in  viscosity  with  slight  change  in 
dilution.  It  is  a  singular  fact  that  Franklin2  found  the  min- 
imum temperature  coefficient  of  conductivity  to  exist  in  the 
solutions  of  intermediate  dilution,  where  the  minimum  con- 
ductivity was  also  found.  This  is  in  accord  with  the  results 
here  given  for  ammonium  thiocyanate  in  pyridine.  On  the 
other  hand,  for  sulphur  dioxide  the  minimum  value  of  the 
equivalent  conductivity  is  displaced  toward  the  region  of 
greater  dilution,  with  rising  temperature,  but  this  is  not  so 
for  solutions  of  the  same  salt  in  pyridine.  It  seems,  there- 
fore, that  the  disturbing  influences  acting  on  ammonium 
thiocyanate  when  dissolved  in  liquid  sulphur  dioxide  are, 
at  least  in  some  measure,  at  work  in  pyridine. 


1  Franklin:  Loc.  cit. 

2  Loc.  cit. 


Electrical  Conductivity  of  Certain  Salts  in  Pyridine      769 

Mercuric  Chloride.  —  The  sample  used  was  recrystallized 
from  conductivity  water  and  thoroughly  dried  at  ioo°-io5°. 

Lang1  and  Reitzenstein2  have  studied  the  compound 
HgCl2  •  Pyr,  and  the  formation  of  HgCl2  •  2Pyr  has  been  in- 
vestigated by  Pesci.3  McBride4  determined  the  tempera- 
ture-solubility curve  for  mercuric  chloride  in  pyridine  at 
temperatures  ranging  from  —  33°  to  +145°  and  proved  the 
existence  of  the  three  compounds,  HgCl2  •  2  Pyr,  HgCl2  •  Pyr 
and  3HgCl2  •  2Pyr. 


XV 


V 

/O 

/2S 

/I0 

0-5 

0.0091 

0.036 

0.045 

I  .0 

0.019 

0.025                0.030 

2  .O 

0.016 

0.021                         0.025 

10.  O 

0.016 

O.O2I                         O.O27 

IOO.O 

0.037 

0.061                  0.067 

1000.  0 

0.130 

0.260                  0.400 

4 

^0° 

/ 

Hc^C 

2   /      y 

W 

2 

/  / 

/  / 

0 

y 

/ 

O 

/, 

/ 

X 

As 

. 

^0* 

1 

sr^ 

jfir 

^^ 

. 

r  Solid 

^ 

^  ^ 

-^ 

\St* 

^ 

-v^ 

0 

v*^1-*^ 

01254 

Fig.  VIII 

1  Ber.  deutsch.  chem.  Ges.,  21,  1578-88  (1888). 

2  Ann.  Phys.  Chem.,  43,  839-40  (1891). 

3  Gazz.  chim.  ital.,  25,  II,  423-33  (1895). 

4  Jour.  Phys.  Chem.,  14,  189-200  (1910). 


770 


Edward  X.  Anderson 
TABLE  XVI 


V 

A 

B 

C 

0.5 

o.  H761 

0.0104 

0.07931 

I  .0 

0.0136 

o  .  0076 

0.0119 

2.0 

0.0126 

o  .  0065 

0.0105 

10.  0 

0.0126 

0.0119? 

0.0141 

IOO.O 

0.0260 

o  .  0039 

0.0162 

IOOO.O 

o  .  0400 

0.0215 

0.0415 

Mercuric  Bromide. — This  salt  was  precipitated  from*  a 
solution  of  mercuric  chloride  with  potassium  bromide,  recrys- 
tallized  from  and  thoroughly  washed  with  conductivity 
water  and  finally  dried  at  ioo°-io5°. 

Groos2  and  Reitzenstein3  made  the  compound  HgBr2  - 
2Pyr. 

TABLE  XVII 


V 

/o 

/25 

fa 

0.5 

O.OI21 

0.034 

0.043 

I  .0 

O.O2O 

O.O26 

0.032 

2.0 

O.OlS 

0.023 

0.026 

IO.O 

0.017 

0.023 

O.O28 

IOO.O 

0.031 

0.047 

0-053 

IOOO.O 

0.13 

O.28                                    O.29 

^-Solid 


HgBr2 


%•- 


Fig.  IX 


1  Solid  phase  present. 

2  Arch.  Pharm.,  (3)  28,  73-8. 
8  Loc.  cit. 


Electrical  Conductivity  of  Certain  Salts  in  Pyridine      771 
TABLE  XVIII 


V 

A 

B 

c 

0.5 

0.07421 

0.0108 

0.05251 

I  .0 
2  .O 
10.  0 

O.OI2I 

0.0104 
0.0128 

o  .  0090 
o  .  0063 
o  .  0092 

0.0119 
0.0092 
0.0124 

100.  0 
IOOO.O 

0.0207 
o  .  0462 

o  .  005  i 
0.0014 

0.0142 
0.0246 

Mercuric  Iodide. — A  method  similar  to  that  used  for  the 
preparation  of  mercuric  bromide  was  followed  in  preparing 
this  salt,  using  potassium  iodide  as  the  precipitant.  Corn- 


X" 

Solid 


50" 


/20° 


0  254 

Fig.  X 

plete^desiccation  was  assumed  after  the  salt  had  been  kept 
for.  some  time  at  the  transition  point  of  the  iodide. 

The  compound  HgI2  •  2Pyr  has  been  prepared  by  Groos.2 
TABLE  XIX 


V 

/o 

b 

b 

0.67 
1  .00 

2  .OO 

Solidified 
0.009 
0.008 

0.013 
0.013 

0.012 

0.018 
0.018 
0.015 

10.00 
100.00 
1000.00 

0.013 
0.069 
0.266 

0.019 
0.  102 
0.364 

0.024 
o.  117 
0.448 

1  Solid  present. 

2  Loc.  cit. 

772 


Edward  X.  Anderson 
TABLE  XX 


V 

A 

B 

C 

0.67 

I  .00 
2.00 
IO.OO 

0.0186 
0.0170 
0.0198 

0.0162 
0.0130 
0.0118 
0.0099 

0.0188 
0.0169 
0.0173 

100.00 
1000.00 

0.0190 
0.0148 

o  .  0059 
o  .  0093 

0.0139 
0.0137 

As  a  general  rule  the  mercuric  halides  when  dissolved 
in  organic  solvents  exhibit  a  strong  tendency  either  to  polym- 
erize, or  to  unite  with  the  solvent  to  form  complex  solvent- 
solute  molecules.  Walden  and  Centnerszwer1  determined 
the  molecular  weights  of  the  three  halides  in  pyridine  by  the 
boiling-point  method  and  found  them  to  be  approximately 
normal  in  the  dilute  solutions.  In  the  concentrated  solu- 
tions, however,  the  molecular  weights  are  less  than  normal, 
which,  since  the  low  conducting  power  is  evidence  of  slight 
ionization,  must  indicate  in  the  solution  the  presence  of  sol- 
vent-solute complex  molecules  and  complex  ions. 

As  was  to  be  expected,  the  equivalent  conductivities 
of  the  three  salts  are  extremely  low.  The  effect  of  tempera- 
ture is  relatively  slight;  the  conductivity  values  for  corre- 
sponding dilutions  are  of  the  same  order  of  magnitude  and 
show  a  minimum  conductivity  in  the  regions  of  greatest 
concentration. 

The  values  of  lv  for  mercuric  iodide  are  much  smaller 
than  those  obtained  by  Lincoln1  at  25°. 

Owing  to  the  low  conductivity  of  these  solutions  and 
the  fact  that  very  slight  errors  are  enormously  magnified 
in  the  calculation,  no  emphasis  is  made  of  the  exactness  of 
the  temperature  coefficients;  they  show,  nevertheless,  dis- 
tinct minima. 

Cupric  Chloride. — This  salt  was  heated  for  several  hours 
in  an  atmosphere  of  dry  hydrogen  chloride  at  i6o°-i65°, 
then  heated  at  the  same  temperature  in  a  current  of  dry 


1  Loc.  cit. 


Electrical  Conductivity  of  Certain  Salts  in  Pyridine      773 

hydrogen  and  cooled  in  a  current  of  the  latter;  lastly,  it  was 
heated  in  an  air  bath  at  about  160°  for  several  hours  and 
preserved  as  above  described. 

Lang1  isolated  the  compound  CuCl2-2Pyr. 

TABLE  XXI 


V 

k 

b 

I* 

25.0 
50.0 

100.  0 
200.0 

500.0 

1000.  0 

0.053 

0.066 

0.088 

o.  130 
0.203 
0.302 

0.062 
0.076 
0.098 
o.  146 
0.216 

0.365 

0.074 
0.086 

0.  Ill 

o.  171 
0.216 
0.410 

TABLE  XXII 

V 

A 

B 

c 

25.0 

o  .  0073 

0.0076 

0.0082 

50.0 

0.0059                0.0052 

0.0059 

IOO.O 
200.0 

0.0045 
0.0050 

0.0055 
0.0068 

o  .  0053 
0.0063 

500.0 

IOOO.O 

0.0027                  o.oooo 
o  .  0084                 o  .  0049 

0.0014 
0.0072 

The  curves  (Fig.  XI)  for  cupric  chloride  show  an  ap- 
parent transition  for  the  values  of  /„.  The  equivalent  con- 
ductivity increases  steadily  with  increasing  dilution.  The 
large  factor  100,  by  which  /  is  multiplied,  is  responsible,  for 
the  marked  break  in  the  curves.  The  temperature  coefficients 
show  minimum  values. 

Ley1  obtained  a  blue  solution  when  he  dissolved  cupric 
chloride  in  pyridine;  this  solution  gave  /30  =  0.05,  in  close 
agreement  with  the  results  given  in  this  work  for  the  same 
dilution.  Ley  assumes  that  the  blue  color  of  the  t  yridine 
solution  is  due  to  an  undecomposed  cupric  chloride-pyridine 
compound  and  not  to  copper  ions.  Kohlschuetter2  states 

1  Loc.  cit. 

2  Ber.  deutsch.  chem.  Ges.,  37,  1153  (1904). 


774 


Edward  X.  Anderson 


that  cupric  chloride  dissolved  in  pyridine  gives  a  blue  solu- 
tion, that  is,  the  color  of  the  solution  corresponds  to  the  color 
of  the  hydrated  salt,  and,  since  its  molecular  weight  as  found 
by  the  boiling-point  method  is  normal,  its  color  may  be  at- 
tributed to  that  of  the  undissociated  cupric  chloride.  In 
these  pyridine  solutions  there  may  be  complexes  of  solvent 
and  solute  of  the  order  of  CuCl2  •  2Pyr,  corresponding  to 
CuCl2  •  2H2O,  or  of  a  higher  order.  Naumann1  states  that 
cupric  chloride  dissolves  in  pyridine  with  evolution  of  heat, 


giving  a  blue  solution  and  therefore  assumes  that  the  com- 
plex CuCl2  •  2Pyr  is  present  in  the  solution.  All  these  inves- 
tigators affirm  that  cupric  chloride  dissolved  in  pyridine 
gives  rise  to  a  blue  solution.  All  of  the  cupric  chloride  solu- 
tions used  in  this  work  had  a  beautiful,  deep  green  color 
without  the  least  indication  of  a  bluish  tint  and,  furthermore, 
the  solutions  remained  green  for  several  months, — until 
finally  rejected.  On  the  other  hand,  in  making  one  of  the 
trial  solubility  tests,  an  attempt  was  made  to  weigh  directly 
a  sample  of  the  dry  cupric  chloride.  The  salt  absorbed 


1  Ber.  deutsch.  chem.  Ges.,  37,  IV,  4609  (1904). 


Electrical  Conductivity  of  Certain  Salts  in  Pyridine      775 

moisture  so  rapidly  that  accurate  weighing  was  impossible. 
Although  it  was  noticed  that  the  edges  of  the  salt  mass  had 
taken  on  a  greenish  blue  color,  the  salt  was  quickly  trans- 
ferred and  dissolved  in  pyridine,  and,  as  might  be  expected, 
the  solution  was  distinctly  blue.  When,  however,  the  cupric 
chloride  was  quickly  weighed  by  difference,  a  deep  green  solu- 
tion was  obtained.  It  is  evident,  therefore,  that  the  blue 
color  observed  by  Kohlschuetter,  Lang,  Ley  and  Naumann 
is  due  to  the  presence  of  a  slight  trace  of  moisture.  It  is 
safe  to  assert  that  the  green  color  of  a  solution  of  cupric 
chloride  in  pyridine  is  due  to  the  presence  of  CuCl2  2Pyr, 
and  the  blue  color  often  obtained  when  supposedly  dry  cupric 
chloride  is  dissolved  in  the  same  solvent  is  due  to  the  pres- 
ence of  CuCl2-2H2O.  The  values  given  by  Lincoln1  at  25° 
for  /„  (recalculated  from  those  given  for  /*)  are  much  higher 
than  the  above  or  those  given  by  Ley.1 

Copper  Nitrate. — The  anhydrous  salt  was  prepared  by 
the  method  of  displacement  used  by  Kahlenberg.2  A  o.i  N 
solution  of  silver  nitrate  in  pyridine  was  treated  with  an  ex- 
cess of  finely  divided,  reduced  metallic  copper  and  allowed 
to  stand  until  the  solution  gave  no  test  for  silver. 

Copper  nitrate  crystallizes  from  pyridine  solutions  as 
the  complex  Cu(NO3)2-4Pyr.3 

TABLE  XXIII 


V 

/. 

/ 

28 

JM 

10.  0 

9.68 

12 

•94 

14.96 

20.0 

5.00 

7 

.21 

8.88 

40.0 

8-57 

ii 

.60 

14.  16 

100.  0 

12.08 

16 

•43 

20.43 

IOOO.O 

16.41 

22 

.88 

29.71 

10,000.0 

19.42 

27 

.24 

35-71 

1  Loc.  cit. 

2  Jour.  Phys.  Chem.,  3,  379  (1899). 

3  Grossmann:  Ber.  deutsch.  chem.  Ges.,  37,  1253-7  (1904). 


Edward  X.  Anderson 


TABLE  XXIV 


V 

A 

B 

C 

IO.O 

0.0135 

o  .  0062 

0.0109 

20.  o 

0.0176 

0.0093   1   0.0155 

40.0 

0.0142 

o  .  0088 

0.0131 

IOO.O 

0.0144 

0.0097 

0.0138 

IOOO.O 

0.0158 

0.0119 

0.0162 

10,000.0 

0.0161 

0.0124       0.0168 

A  study  of  Fig.  XII  discloses  the  fact  that  copper  nitrate 
is  peculiar  in  its  behavior,  in  that,  like  the  mercuric  halides, 
it  gives  minimum  conductivity  values.  After  passing  through 
the  minima  the  values  for  lv  first  increase  at  a  moderate  rate 
and  then  exceedingly  slowly  with  increasing  dilution. 


Here  again,  the  temperature  coefficients  show  minimum 
values. 

Cadmium  Nitrate. — The  pure  anhydrous  salt  was  pre- 
pared by  the  displacement  of  silver  in  a  o.i  N  solution  of 
silver  nitrate  by  means  of  chemically  pure  cadmium. 

TABLE  XXV 


V 

/O                                                /,5 

ho 

IO.O 

o.  141 

o.  1  60 

O.  122 

20.0 

0.322 

0.348 

0.288 

40.0 

0.402 

0-433 

0.340 

IOO.O 

0.694 

0-733 

0.630 

IOOO.O 

2.37 

2.31 

2.44 

10,000.0 

7.40 

8.60 

9.80 

Electrical  Conductivity  of  Certain  Salts  in  Pyridine      777 
TABLE  XXVI 


V 

A 

B 

c 

10.  0 

0.0052 

-0.0095 

—  0.0028 

20.0 

0.0033 

—0.0070 

—  0.0021 

40.0 

0.0031 

-^0.0086 

—0.0031 

100.0        0.0023 

—0.0056 

—0.0018 

1000.0      —  o.ooio 

o  .  0023 

0.0006 

10,000.0        0.0065 

0.0056 

0.0065 

The  equivalent  conductivities  of  cadmium  nitrate  (Fig. 
XIII)  increase  quite  steadily  with  increase  in  dilution.  The 
curves  obtained  by  plotting  the  temperature  coefficients 
are  somewhat  irregular,  but  in  general  show  minima  of  nega- 
tive value.  The  curves  run  quite  closely  together,  diverging 
only  to  an  appreciable  extent  in  the  more  dilute  solutions. 


4 

Fig.  XIII 

Cobalt  Chloride. — The  sample  used  was  partially  dehy- 
drated in  vacuo  over  phosphorus  pentoxide,  then  heated 
in  an  atmosphere  of  dry  hydrogen  chloride  at  140°  for  twenty- 
four  hours,  and  finally  in  a  current  of  dry  hydrogen  for  fif- 
teen hours.  The  product  was  of  a  pale  blue  color.  Reitzen- 
stein1  prepared  the  compound  CoCl2  •  4Pyr.  Cobalt  chloride 

1  Loc.  cit. 


778 


Edward  X.  Anderson 


dissolved  in  pyridine  gives  a  red  solution  at  o°,  a  violet  at 
25°,  and  a  deep  purple  solution  at  50°. 

TABLE  XXVII 


V 

/o 

fa 

fa 

IO.O 

0.0091 

0.012 

0.012 

20.  o 

0.015 

0.015 

O.O22 

40.0 

0.021 

O.O2O 

0.024 

100.  0 

0.042 

0.045 

0.041 

IOOO.O 

O.22O 

0.230 

0.310 

10,000.0                    0.600 

I  .OOO                            I  .  IOO 

TABLE  XXVIII 

V 

A 

B                                   C 

IO.O 

O.OI491 

0.0319                     0.02931 

20.0 

0.0019 

0.0174                     o.oioi 

40.0 

—  o  .  002  i 

0.0082                    0.0028 

100.  0 

0.0028 

—  0.0036                   —  0.0005 

IOOO.O                           O.OOlS 

0.0139                       0.0082 

10,000.0                  0.0267 

0.0040                      0.0167 

6 

L/V 

yl 

/| 

/ 

/    Cod2 

'          / 

/  25 

4 

/ 

/ 

^^- 

2 

/ 

// 

^ 

J^ 

§ 

/^ 

7^ 

X 

/ 

~? 

X 

/^ 

\ 

•Ss^ 

o 

[2^M<r 

^ 

v> 

0            1            2            3>           4 

678 

Fig.  XIV 

Solid  present. 


Electrical  Conductivity  of  Certain  Salts  in  Pyridine      779 

Cobalt  chloride  in  pyridine  solutions  is  at  best  an  ex- 
ceedingly poor  conductor.  By  some,  its  solutions  are  con- 
sidered as  non-conductors.  Consequently,  slight  errors  in 
the  work  are  highly  magnified.  The  results  obtained  show 
a  continuous  increase  in  equivalent  conductivity  with  dilu- 
tion for  all  temperatures.  Again,  Lincoln's1  values  for  /„ 
at  corresponding  dilutions  are  very  much  higher  than  the 
values  here  quoted.  The  temperature  coefficients  are  like- 
wise subject  to  considerable  error,  yet  even  these  show  definite 
minima  at  which  negative  coefficients  are  observed. 


Fig.  XV 

Werner1  and  his  co-workers  found  by  the  ebullioscopic 
method  that  cobalt  chloride  has  the  normal  molecular  weight 
in  pyridine. 

t£  Lead  Chloride. — The  salt  was  precipitated  from  chemically 
pure  lead  nitrate  by  pure  hydrochloric  acid.  It  was  then 
thoroughly  washed  with  water  and  heated  to  dryness  at  120°. 

1  Loc.  cit. 


Edward  X.  Anderson 


Pyridine  forms  three  crystalline  compounds  with  lead 
chloride,  e.  g.,  3PbCl2  4Pyr,1  4PbCl2  •  3?yr,2  and  PbQ2  2Pyr.3: 
It  crystallizes  from  aqueous  solutions  in  the  form  of  the  an- 
hydrous salt  and  should  have  small  hydrating  power  and 
low  temperature  coefficients  of  conductivity.  Jones  and 
Winston4  find,  however,  that  the  temperature  coefficients  in 
aqueous  solutions  are  relatively  high.  Apparently  the  tem- 
perature coefficients  are  not  entirely  dependent  upon  the 
amount  of  solvent  separating  with  the  solute  when  solvates 
crystallize  from  solution. 

TABLE  XXIX 


V 

/o 

h 

/IO 

50.0 

0.05 

0.07 

0.  10 

IOO.O 

O.O6 

0.  10 

0.15? 

200.0 

0.07 

O.  IO 

O.  II 

4OO.O 

0.  12 

0.18 

0.23 

8OO.O 

O.2O 

0.26 

0.35 

1600.  o 

0-35 

0.48                      0.63 

TABLE  XXX 

V 

A 

B                                 C 

50.0 

0.030 

O.OIQ                               O.O2O 

IOO.O 

0.025 

0.021                               0.030? 

200.0 

0.019 

0.005?                            0.013 

400.0 

0.018 

0.013                     0.018 

800.0 

0.014 

O.OI2                                O.OI5 

1600.  o 

0.016 

0.012                               0.016 

Discussion 

The  power  of  a  solution  to  conduct  electricity  depends 
upon  the  dielectric  constant  of  the  solvent,  the  degree  of  dis- 
sociation of  the  electrolyte,  the  number  of  ions  present  and 
their  velocities.  The  degree  of  dissociation  is  determined 


1  Zeit.  anorg.  Chem.,  4,  100-110  (1893). 

2  Ibid.,  14,  379-403  (1897)- 

3  Jour.  Phys.  Chem.,  15,  373  (1912). 

4  Am.  Chem.  Jour.,  46,  368  (1912). 


Electrical  Conductivity  of  Certain  Salts  in  Pyridine      781 

by  the  dielectric  constant  of  the  solvent  and  the  electro- 
affinity  of  the  ions.  For  a  given  potential  gradient  the  ionic 
velocities  are  in  turn  dependent  upon  the  viscosity  of  the 
solution,  the  mass  or  the  volume  of  the  ions.  The  mass  and 
volume  of  either  or  both  ions  may  be  further  augmented 
by  combination  with  the  molecules  of  the  solute  to  form 
complex  ions,  or  with  the  molecules  of  the  solvent  to  form 
more  or  less  highly  solvated  ions. 

For  solutions  in  a  large  number  of  solvents  the  molecular 
conductivity  behaves  normally,  i.  e.,  it  increases  with  in- 
creasing dilution.  There  are,  however,  numerous  instances 
like  those  shown  by  the  mercuric  halides  in  pyridine  in  which 
the  equivalent  conductivity  decreases,  passes  through  a 
minimum  and  then  rises  with  increase  in  the  concentration 
of  the  solute.  Such  phenomena  are  considered  by  many 
as  being  incapable  of  explanation  on  the  basis  of  the  Arrhenius 
theory.  That  increase  in  molecular  conductivity  with  in- 
crease in  concentration  in  some  organic  solvents  is  peculiar 
only  to  the  concentrated  solutions  has  been  shown  by  Pearce1 
and  others.  He,  like  Sachanov,2  found  that  the  molecular 
conductivity  of  various  salts  in  aniline  decreased  with  in- 
crease in  dilution  to  a  minimum  and  then  increased  normally 
upon  further  dilution. 

Franklin3  explains  the  increase  in  molecular  conductivity 
with  increasing  concentration  as  due  to  an  increase  in  auto- 
ionization,  which  more  than  compensates  for  the  decrease  in 
molecular  conductivity  due  to  an  increase  in  viscosity. 

In  their  study  of  the  conductivity  of  various  inorganic 
salts  and  non-salt  organic  solutes  in  the  liquid  halogen  acids 
Archibald,  Mclntosh  and  Steele4  found  that  the  molecular 
conductivity  increases  with  increasing  concentration.  They 
assumed  that,  while  the  original  solute  may  of  itself  be  in- 
capable of  ionizing  in  the  solvent,  it  may  combine  with  the 

1  Unpublished  results. 

2  Jour.  Russ.  phys.  chem.  Soc.,  42,  683-690  (1910). 

3  Loc.  cit. 

4  Zeit.  phys.  Chem.,  55,  129  (1907)- 


782  Edward  X.  Anderson 

solvent  to  form  a  complex  solvent-solute  molecule  which 
behaves  as  an  electrolyte.  Accordingly,  the  number  of  com- 
plexes must  increase  with  the  concentration  and,  likewise, 
the  molecular  conductivity  must  increase,  due  to  an  increase 
in  the  total  number  of  ions.  These  investigators  used  the 
expression 

Mv  =  aK'  =  XVn, 

where  a,  x  and  V  represent  the  degree  of  dissociation,  the 
specific  conductance  and  the  volume,  respectively.  K'  is 
a  constant  and  n  is  the  number  of  simple  molecules  of  solute 
combined  with  m  molecules  of  solvent.  By  means  of  this 
relation  they  calculated  the  molecular  conductivity  and 
found  it  to  increase  normally  with  dilution.  It  is  obvious 
that  when  a  is  equal  to  unity, 

Moo   =  K'. 

Evidently  these  investigators  did  not  consider  the  fact 
that  the  amount  of  solvation  per  molecule  of  solute  will 
vary  with  the  dilution  and  must  necessarily  decrease  as  the 
concentration  of  the  solute  is  increased.  Where  solvation 
is  possible  it  is  difficult  to  conceive  how  the  complex,  say, 
(n)  solute-  (m)  solvent  can  remain  unchanged  throughout  a 
wide  range  in  concentration. 

Sachanov  has  studied  the  molecular  conductivity  of 
various  solutions  in  acetic  and  propionic  acids1  and  in  aniline,, 
methylaniline  and  dimethylaniline.2  In  practically  every 
case  the  molecular  conductivity  was  found  to  increase  with 
increasing  concentration.  All  of  these  solvents  have  low 
dielectric  constants  and  slight  dissociating  power.  In  a  later 
article,3  he  states  that  a  decrease  in  molecular  conductivity 
with  increasing  dilution  is  just  as  characteristic  for  solvents 
with  low  dielectric  constants  as  is  an  increase  in  molecular 
conductivity  for  solutions  in  solvents  with  high  dielectric 
constants.  He4  asserts  that  electrolytic  dissociation  does 

1  Jour.  Russ.  phys.  chem.  Soc.,  43,  526  (1911). 

2  Ibid.,  44,  324  (1912);  42,  683  (1910). 

3  Zeit.  phys.  Chem.,  80,  13  (1912). 

4  Ibid.,  80,  20  (1912). 


Electrical  Conductivity  of  Certain  Salts  in  Pyridine      783. 

not  depend  solely  on  the  magnitude  of  the  dielectric  constant, 
but  also  on  the  solvates  and  complex  ions.  The  formation 
of  such  ions  favors  electrolytic  dissociation  because  the  elec- 
tro-affinity of  these  complex  ions  is  greater  than  that  of  the 
simple  ions.  In  solvents  with  low  dielectric  constants  only 
the  complexes  which  yield  complex  ions  of  high  electro- 
affinity  can  undergo  electrolytic  dissociation.  The  decrease 
in  electrolytic  conductivity  in  such  solvents  is  explained  as 
being  due  to  the  composition  of  these  polymerized  solute 
molecules  on  dilution. 

An  examination  of  the  curves  plotted  from  the  data 
will  show  that  the  electrolytes  used  in  this  work  are  of  two 
types.  First,  those  which  give  minimum  conductivity  values 
at  all  temperatures,  viz.,  the  mercuric  halides  and  copper 
nitrate.  Second,  those  like  the  nitrates  of  silver  and  cadmium, 
the  chlorides  of  lithium,  cobalt,  copper  and  lead,  the  iodides 
of  lithium  and  sodium,  lithium  bromide,  and  the  thiocyanates 
of  potassium  and  ammonium,  which  upon  dilution  give  in- 
creasing values  of  /„. 

The  phenomenon  of  the  molecular  conductivity  increas- 
ing with  increasing  concentration  is  considered  by  some  as 
being  at  variance  with  the  Arrhenius  theory.  However, 
an  attempt  will  be  made  in  the  present  discussion  to  show  that 
this  anomalous  behavior  is  due  entirely  to  the  presence  and 
the  properties  of  the  polymerized  solute  molecules  which 
are  present  in  solutions  showing  such  phenomena. 

The  most  dilute  solutions  studied  by  Walden  and  Cent- 
nerszwer1  correspond  to  those  concentrations  which  give 
the  minimum  conductivity  values  in  this  work,  and  hence 
we  have  no  experimental  evidence  as  to  the  state  of  aggrega- 
tion in  the  very  dilute  solutions.  Assuming  only  the  ioniza- 
tion  of  simple  molecules,  their  results  would  seem  to  show 
that  these  molecules  are  ionized  in  the  concentrated,  but 
undissociated  in  the  dilute  solutions,  which  is  contrary  to 
common  experience.  That  some  ions  are  present  at  all  dilu- 
tions is  evident  from  the  conducting  power  of  tlu>r  solutions. 

1  Zcit.  phys.  Chem.,  55,  231    (1906). 


784  Edward  X.  Anderson 

The  abnormal  boiling-point  results  noted  must,  therefore, 
be  due  to  such  an  equilibrium  between  the  various  forms  of 
the  electrolyte  present  in  the  solution,  viz.,  simple  and  com- 
plex molecules,  simple  and  complex  ions  and  the  pyridinated 
forms  of  each,  as  would  give  the  molecular  weights  obtained. 

Let  us  consider,  as  an  example  of  the  first  type,  the  solu- 
tions of  mercuric  chloride,  and  what  is  said  concerning  them  will 
apply,  for  the  most  part,  to  the  other  salts  which  give  min- 
imum conductivity.  This  salt  is  very  soluble  and  has  a  strong 
tendency  to  combine  with  pyridine  to  form  stable  salts  with 
pyridine  of  crystallization.  Naturally,  then,  we  would  ex- 
pect it  to  combine  with  still  more  pyridine  when  in  solution 
in  this  solvent.  Walden  and  Centnerszwer1  have  found 
that  the  mercuric  halides  exhibit  a  strong  tendency  to  polym- 
erize in  pyridine  solutions,  just  as  they  do  in  most  solvents. 

We  should  expect  to  find  in  these  solutions  the  simple 
and  polymerized  molecules  of  mercuric  chloride,  the  simple 
and  complex  ions  and  the  solvated  ions  and  molecules.  The 
number  and  kind  of  each  will  depend  upon  the  dilution. 
The  equilibria  between  the  different  forms  of  the  solute 
may  be  represented  thus: 

(1)  HgCl2  -f  WiPyr  ±Z^  HgCl2.WiPyr 

(2)  2HgCl2  -  mtPyr  ±^  (HgCl2)2  .  w2Pyr 

As  the  concentration  of  the  solute  is  increased  the  equi- 
librium will  naturally  be  disturbed  and  the  resulting  effect 
will  be  a  corresponding  increase  in  the  concentration  of  the 
polymer.  Hence  the  reaction  will  proceed  from  left  to  right, 
as  in  (2).  It  is  evident  from  the  increase  in  conductivity 
in  the  most  concentrated  solutions  that  the  dissociation  of 
the  polymerized  molecules  must  be  greater  than  that  of  the 
simple  molecules.  The  conditions  of  equilibria  may  be  repre- 
sented by  the  following  relations: 


(3)    (HgCl2)2  -  w2Pyr  ±^:  Hg  •  w3Pyr  +  HgCU  '  w4Pyr 

In  the  most  dilute  solutions  where  the  simple  molecules 
predominate,  we  shall  have  the  equilibrium  represented  by 


1  Loc.  cit. 


Electrical  Conductivity  of  Certain  Salts  in  Pyridine      785 

(4)     HgCl2  -  WiPyr  ±^:  HgCl  -  w5Pyr  +  Cl .  w6Pyr, 

and  possibly  to  a  slight  extent, 

+  ++ 

(5)     HgCl-w5Pyr  ^:  Hg-w3Pyr  +  Cl  m6Pyr. 

In  the  intermediate  concentrations  all  of  these  equilibria 
will  be  found  to  a  greater  or  lesser  extent.  The  degree  of  pyri- 
dination  of  the  various  solute  particles  (mi,  w2,  etc.)  will 
depend  upon  the  dilution. 

Inasmuch  as  the  molecular  weight  of  mercuric  chloride 
in  pyridine  approximates  the  normal  value,  and,  therefore, 
the  effect  due  to  polymerization  must  approximately  annul 
that  due  to  dissociation  and  solvation,  the  extent  of  the  poly- 
merization must  be  slight,  because  the  conductivity  is  low. 
Hence  we  may  attribute  to  the  polymerized  molecules  a  strong 
ionizing  power  and,  consequently,  a  high  degree  of  dissocia- 
tion in  these  solutions. 

All  the  solutions  used  in  the  present  work  were  made  up 
under  the  assumption  that  no  polymerization  of  the  solute 
takes  place.  It  is,  therefore,  evident  that  the  actual  normali- 
ties of  the  concentrated  solutions  with  respect  to  the  ionizing 
polymers  are  less  than  assumed  in  preparing  the  solutions. 
For  example,  if  exactly  one  gram-molecular  weight  of  mer- 
curic chloride  is  dissolved  in  pyridine  and  then  diluted  to 
one  liter  and,  if  it  is  assumed  that  all  the  simple  molecules 
combine  in  pairs  to  form  single  molecules  of  the  polymerized 
form,  the  normality  with  respect  to  the  polymer  will  be  o .  5  N 
instead  of  i .  o  N,  as  would  be  the  case,  if  no  polymerization 
takes  place.  If,  on  the  other  hand,  only  a  fraction  of  the 
simple  molecules  combine  to  form  the  polymerized  form, 
the  normality  of  the  ionizing  polymers  will  be  still  less,  and 
will  always  be  one-half  the  fractional  molar  concentration 
of  the  simple  molecules  which  have  become  polymerized. 
If  we  are  then  dealing  with  ionizing  polymers,  the  degree  of 
dissociation  for  a  given  solution  in  the  concentrated  regions 
will  be  relatively  much  higher  than  it  would  be  if  we  were 
actually  dealing  with  the  normality  at  which  the  solutions 
are  originally  made  up. 


786  Edward  X.  Anderson 

In  order  to  explain  minimum  conductivity  the  following 
equilibrium  equation  will  be  used  as  a  guide: 

2Hg++  +  4C1-  ±^:  2HgCi+  +  2d-  ±^: 

2HgCl2  ±^:  (HgCl2)2  ±^:  Hg++  +  HgCl4- 

In  the  concentrated  solutions  the  ions  present  will  re- 
sult from  the  dissociation  of  the  polymers.  There  will  be 
only  a  very  slight  tendency  towards  simple  ionization  and 
this  will  be  impeded  by  the  repression  of  ionization  due  to 
the  mercury  ions.  As  the  most  concentrated  solutions  are 
dilute  the  equilibrium  will  be  disturbed  and  polymers  will 
change  over  into  simple  molecules  in  obeyance  to  the  law 
of  mass  action.  Hence  the  concentration  of  the  ions  formed 
from  the  polymers  will  decrease,  although  the  dissociation 
of  the  remaining  polymers  will  become  more  and  more  com- 
plete with  continued  dilution.  The  result  will  be  a  decrease 
in  the  equivalent  conductivity.  As  dilution  is  continued 
the  influence  of  the  simple  ionization  will  soon  become  effec- 
tive on  account  of  the  polymers  being  converted  almost  com- 
pletely into  simple  molecules.  These  in  turn  furnish  a  pre- 
ponderance of  simple  ions  which  will  cause,  the  repression 
of  the  complex  ionization.  The  equivalent  conductivity 
should,  therefore,  begin  to  rise  again  and  continue  to  do  so 
with  an  increase  in  dilution.  This,  then,  would  mean  that 
the  minimum  values  would  be  encountered  for  the  equivalent 
conductivity.  Of  course  viscosity  changes  as  well  as  solva- 
tion  effects  are  taking  place  with  dilution.  The  latter  would 
naturally  change  the  sizes  of  the  ions  and,  therefore,  alter 
the  freedom  with  which  they  would  move.  The  decrease  in 
viscosity  would  also  result  in  greater  freedom  to  the  mobility 
of  the  ions.  However,  all  these  last  mentioned  factors  would 
be  subordinated  to  the  changes  in  the  concentrations  of  the 
various  kinds  of  ions,  as  just  outlined. 

Solutions  of  silver  nitrate,  lithium  bromide,  lithium 
iodide,  and  the  thiocyanates  of  ammonium  and  potassium 
all  possess  high  conductivity  values.  The  most  concentrated 
solution  of  each  is  noticeably  viscous.  On  the  other  hand, 


Electrical  Conductivity  of  Certain  Salts  in  Pyridine      787 

cadmium  nitrate  and  the  chlorides  of  lithium,  cobalt,  copper 
and  lead  are  only  slightly  soluble  in  pyridine.  Their  most 
concentrated  solutions  have  apparently  low  viscosities.  The 
equivalent  conductivities  shown  by  these  electrolytes  are  very 
low.  Hence,  within  the  limits  of  these  facts,  the  viscosity 
is  not  the  predominating  effect  in  their  conductivity. 

Another  point  in  connection  with  the  salts  just  men- 
tioned is  the  fact  that,  generally  speaking,  the  binary  salts 
are  the  better  and  the  ternary  salts  the  poorer  conductors. 

The  existence  of  minima  of  temperature  coefficients 
of  equivalent  conductivity  seems  to  be  a  general  phenomenon 
in  pyridine  solutions.  The  solvation  per  molecular  weight 
of  the  solute  is  greatest  in  the  more  dilute  solutions.  Owing 
to  the  greater  concentrations  of  the  solute,  however,  the  con- 
centrated solutions  will  contain  a  much  greater  percent  of 
combined  pyridine,  the  degree  of  solvation  decreasing  with 
increasing  concentration.  The  instability  of  these  solvates 
increases  not  only  with  increase  in  complexity  but  also  with 
increase  in  temperature. 

The  minima  of  the  temperature  coefficients  may  be  ex- 
plained on  the  basis  of  viscosity  and  ionic  solvation.  We 
would  expect  the  greatest  temperature  changes  in  the  most 
concentrated  solutions.  Here  the  viscosity  is  greatest  and 
the  percent  of  pyridine  locked  up  as  that  of  solvation  is  at  a 
maximum.  The  effect  of  the  decrease  in  the  complexity  of 
the  solvates  due  to  a  rise  in  temperature  would  be  to  increase 
the  amount  of  the  pure  solvent.  This  in  itself  means  a  de- 
crease in  the  viscosity  of  the  solution.  Then,  too,  the  rise 
in  temperature  results  in  a  decrease  in  the  viscosity  of  the 
pure  solvent.  For  these  reasons  the  solvated  ions,  which  in 
themselves  may  or  may  not  have  suffered  an  appreciable 
change,  due  to  the  rise  in  temperature,  are  able  to  migrate 
more  rapidly.  In  the  less  concentrated  solutions  the  vis- 
cosity changes  due  to  a  rise  in  temperature  will  be  less  pro- 
nounced. Hence  the  values  of  the  temperature  coefficients 
will  decrease.  While  the  decrease  in  the  viscosity  change 


788  Edward  X.  Anderson 

with  dilution  is  going  on  the  ionic  solvation  is  increasing. 
The  instability  of  the  solvated  ions  with  rise  in  temperature 
is  also  increasing  with  dilution.  As  the  temperature  is  raised 
the  decrease  in  the  complexity  of  the  solvated  ions  will  cause 
an  increase  in  their  mobility.  Thus  we  can  readily  under- 
stand how  the  solvated-ion  effect  can  annul  that  of  the  vis- 
cosity. It  is  this  condition  which  gives  rise  to  the  minimum 
temperature  coefficient  values.  As  the  increase  in  dilution 
continues  the  ion  solvation  effect  becomes  the  sole  factor 
and  hence  we  obtain  a  steady  increase  in  the  values  of  the 
temperature  coefficients. 

Another  point  to  be  noted  is  that  the  temperature  co- 
efficients are  greater  between  o°  and  25°  than  those  obtained 
for  the  interval  25°-5o°.  Since  the  decrease  in  the  com- 
plexity of  the  solvated  ions  caused  by  a  rise  in  temperature 
must  be  greatest  in  the  region  of  greatest  complexity,  a  greater 
change  should  be  expected  at  temperatures  favorable  to  a 
higher  degree  of  complexity,  i.  e.,  between  o°  and  25°. 

No  explanation  can  be  offered  for  the  negative  tempera- 
ture coefficients  in  the  cases  of  lithium  bromide,  sodium 
iodide,  cobalt  chloride  and  cadmium  nitrate. 

Comparing  the  magnitudes  as  well  as  the  relative  changes 
of  temperature  coefficients  with  dilution  for  the  different 
electrolytes  there  seems  to  be  no  direct  relation  between 
these  and  the  amount  of  pyridine  with  which  they  are  com- 
bined when  they  crystallize  from  solution. 

In  general,  the  binary  salts  show  the  highest  equivalent 
conductivity.  The  ternary  give  the  lowest.  There  are, 
however,  two  notable  exceptions  to  this  generalization: 
lithium  chloride,  a  binary  salt,  gives  very  low,  while  copper 
nitrate,  a  ternary  salt,  gives  comparatively  high  values  of  the 
equivalent  conductivity. 

So  far  as  the  effect  of  sedation  on  conductivity  is  con- 
cerned, there  is  no  reason  for  assuming  the  combination  of 
solute  and  solvent  as  essential  to  the  increase  of  electro- 
affinity,  as  Sachanov1  has  done.  After  the  ions  are  once 

1  Loc.  cit. 


Electrical  Conductivity  of  Certain  Salts  in  Pyridine      789 

formed  then  combination  of  the  same  with  the  solvent  may 
modify  their  electro-affinity  to  some  extent. 

Inasmuch  as  the  dielectric  constant  of  pyridine  (20) 
varies  little  within  the  temperature  range  studied  its  effect 
must  be  practically  constant  throughout.  Since  pyridine 
is  considered  as  a  non-associated  solvent1  this  factor  would 
not  enter  into  the  problem. 

Summary 

A  study  of  the  equivalent  electrical  conductivity  of  solu- 
tions of  fifteen  salts  in  pyridine  has  been  made  at  three  tem- 
peratures, viz.,  o°,  25°  and  50°. 

Two  classes  of  electrolytes  are  to  be  observed,  (i)  Those 
for  which  the  equivalent  conductivity  increases  throughout 
with  increasing  dilution;  (2)  those  which  give  minimum 
values  of  equivalent  conductivity. 

The  values  of  the  temperature  coefficients  between  o° 
and  25°  are  higher  than  those  between  25°  and  50°.  An 
explanation  is  offered. 

An  explanation  for  the  minimum  values  of  equivalent 
conductivity  has  been  advanced. 

It  has  been  shown  that  the  anomalous  behavior  in  equiva- 
lent conductivity  is  due  entirely  to  the  presence  and  proper- 
ties of  the  ionizable  polymerized  solute,  which  predominates 
in  the  concentrated  solutions. 


1  J.  L.  R.  Morgan:  Jour.  Am.  Chem.  Soc.,  30,  1068  (1908). 


BIOGRAPHY. 

The  author  was  born  in  Minneapolis,  Minnesota,  February 
9,  1885,  and  began  his  school  career  in  the  public  schools  of 
the  same  city,  graduating  from  the  eighth  grade  in  1899.  He 
next  attended  the  East  Minneapolis  High  School,  receiving  a 
diploma  in  1903.  A  course  in  Analytical  Chemistry  was  then 
pursued  at  the  University  of  Minnesota,  and  the  degree  B.S. 
in  Chemistry  was  received  in  1908.  The  following  year  the 
degree  M.S.  was  conferred  upon  him  at  the  same  institution. 
While  at  the  University  of  Minnesota  he  assisted  in  Qualita- 
tive Analysis  for  five  years.  From  1909  to  1913  he  held  a 
position  as  Instructor  in  Chemistry  at  the  State  University  of 
Iowa,  and  since  then  has  been  instructing  in  Chemistry  at  the 
State  University  of. North  Dakota. 


ACKNOWLEDGMENT. 

The  following  investigation  was  made  at  the  suggestion  of 
Dr.  J.  N.  Pearce,  and  followed  out  under  his  direct  supervision. 
The  author  takes  this  opportunity  to  express  his  sincere  thanks 
to  Dr.  Pearce  for  his  kindly  interest  and  valuable  assistance 
throughout  the  work. 

The  author  also  extends  thanks  to  Dr.  E.  W.  Rockwood 
and  Dr.  W.  J.  Karlslake  for  constant  advice  and  suggestions. 

E.  X.  A. 


A* 


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